Properties, Reactions, and Calculating pH Acids and Bases Properties, Reactions, and Calculating pH
Standards **STANDARD SET 5: Acids and Bases 5. Acids, bases, and salts are three classes of compounds that form ions in water solutions. As a basis for understanding this concept: a. Students know the observable properties of acids, bases, and salt solutions. 5. b. Students know acids are hydrogen-ion-donating and bases are hydrogen-ion-accepting substances. 5. c. Students know strong acids and bases fully dissociate and weak acids and bases partially dissociate. 5. d. Students know how to use the pH scale to characterize acid and base solutions. 5. e.* Students know the Arrhenius, Brønsted-Lowry, and Lewis acid–base definitions. 5. f.* Students know how to calculate pH from the hydrogen-ion concentration. 5. g.* Students know buffers stabilize pH in acid–base reactions.
Properties Acids Electrolytes React with metals to produce H2 gas Form H3O+ ions in water Tastes sour Examples: vinegar, lemon juice, battery acid Bases slippery Form OH– ions in water Tastes bitter Examples: soap, cocoa powder, ammonia, Drano
Calculating pH [H3O+] = 10–pH H+ concentration= 1 × 10–12 pH = 12
pH < 7 = acidic pH > 7 = basic pH Scale pH = 7 = neutral pH 1 14 Strong acids Weak acids Neutral : water Weak bases Strong bases pH < 7 = acidic pH = 7 = neutral pH > 7 = basic
Various Definitions of Acids/Bases Arrhenius Acid – a substance which forms hydronium (H3O+) ions in water. Arrhenius Base – a substance which forms hydroxide (OH–) ions in water. Brønsted-Lowry Acid – a substance which can donate a proton (an H+). Brønsted-Lowry Base – a substance which can accept a proton (an H+). Lewis Acid – a substance which can accept a lone pair of electrons. Lewis Base – a substance which can donate a lone pair of electrons.
Acid Reactions Strong acid: HCl (aq) + H2O (l) Cl– (aq)+ H3O+ (aq) hydronium ion Weak acid: CH3COOH(aq) +H2O(l) CH3COO–(aq)+ H3O+(aq)
Base Reactions Strong base: NaOH (aq) + H2O (l) Na+ (aq)+ H2O + OH–(aq) hydroxide ion Weak base: NH3 (aq) + H2O(l) NH4+(aq)+ OH–(aq)
Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l) CH3COO–(aq)+ H3O+(aq) conjugate base acid NH3 (aq) + H2O(l) NH4+(aq)+ OH–(aq) base conjugate acid
Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l) CH3COO–(aq)+ H3O+(aq) conjugate base acid NH3 (aq) + H2O(l) NH4+(aq)+ OH–(aq) base conjugate acid
Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l) CH3COO–(aq)+ H3O+(aq) H+ conjugate base acid H+
Conjugate Acid/Base Pairs NH3 (aq) + H2O(l) NH4+(aq)+ OH–(aq) H+ H+ base conjugate acid
Water is an Acid/Base H2O (l) + H2O (l) H3O+ (aq) + OH–(aq) hydronium ion hydroxide ion But water is a weak acid and a weak base, so only a small amount of water will form these ions, the rest will remain as H2O.
Measuring the Strength of Acids/Bases We often speak of acid concentrations in molarity (ex. 2.0 M HCl solution, aka ) 2.0 mol HCl 1 L solution But which is stronger 2.0 M HCl solution or 2.0 M CH3COOH solution? It’s more important to know the concentration of hydronium ions. [H+] really means [H3O+]
10 10 [H3O+] = 10–pH Calculating pH pH = – log10[H3O+] (–1) x Example: If [H3O+] = 1 x 10–5 M, what is the pH? pH = – log10[1 x 10–5] pH = 5 Back-calculating concentration (–1) x – log10[H3O+] = pH x (–1) log10[H3O+] = –pH 10 10 [H3O+] = 10–pH
H2O (l) + H2O (l) H3O+ (aq) + OH–(aq) hydronium ion hydroxide ion
H3O+ (aq) + OH–(aq) H2O (l) + H2O (l)
H2O (l) + H2O (l) H3O+ (aq) + OH–(aq)
H3O+ (aq) + OH–(aq) H2O (l) + H2O (l) neutral solution
H3O+ (aq) + OH–(aq) H2O (l) + H2O (l) acidic solution
H3O+ (aq) + OH–(aq) H2O (l) + H2O (l) basic solution
Calculating [OH–] from [H3O+] H2O (l) + H2O (l) H3O+ (aq) + OH–(aq) Kw = [H3O+]·[OH–] Kw = 1 x 10–14 always this number Example: If [H3O+]= 1 x 10–5M, what is [OH–]? 1 x 10–14 = [1 x 10–5M]·[OH–] [1 x 10–5M] [1 x 10–5M] 1 x 10–9 M = [OH–]
pH < 7 = acidic pH > 7 = basic pH Scale pH = 7 = neutral pH 1 14 Strong acids Weak acids Neutral : water Weak bases Strong bases pH < 7 = acidic pH = 7 = neutral pH > 7 = basic
Neutralization H3O+ (aq) + OH–(aq) H2O (l) + H2O (l) Large amounts of H3O+ and OH– cannot exist at the same time in a solution. Whichever ion has the larger amount will reduce the lesser amount. What you would see: HCl (aq) + NaOH (aq) NaCl (aq)+ H2O (l)
Titration H3O+ (aq) + OH–(aq) H2O (l) + H2O (l) Titration is the experimental process of figuring out the pH of a mystery solution by neutralizing it by incrementally adding small amounts of a known solution. The titration is complete when we reach the equivalence point (where [H3O+]=[OH–] ). n = C x V moles = concentration x volume
Titration H3O+ (aq) + OH–(aq) H2O (l) + H2O (l) Example: How many moles of H3O+ would it take to neutralize 5 liters of a 0.1 M NaOH solution? n = C x V n = (0.1 )·(5 L) mol L n = 0.5 mol H3O+
Buffers Buffers stabilize the pH of acids. The salt of the conjugate base to the acidic solution can act as a buffer. CH3COOH(aq)+H2O(l) CH3COO–(aq)+ H3O+(aq) conjugate base acid How would the equilibrium shift if we added NaCH3COO to the acidic solution? CH3COO– adding more
He H B C N O Ne F Li Be P Al Si S Cl Ar Na Mg Br Kr K Ca I Xe
He H B C N O Ne F Li Be P Al Si S Cl Ar Na Mg Br Kr K Ca I Xe
4 e– in valence shell
Measuring the Strength of Acids/Bases We often speak of acid concentrations in molarity (ex. 2.0 M HCl solution, aka ) 2.0 mol HCl 1 L solution But which is stronger 2.0 M HCl solution or 2.0 M CH3COOH solution? Since we use both strong and weak acids, a more consistent measurement would tell us just the concentration of hydronium ions. [H+] really means [H3O+]
H3O+ (aq) + OH–(aq) H2O (l) + H2O (l)
H3O+ (aq) + OH–(aq) H2O (l) + H2O (l)
Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l) CH3COO–(aq)+ H3O+(aq) conjugate base acid NH3 (aq) + H2O(l) NH4+(aq)+ OH–(aq) base conjugate acid
Conjugate Acid/Base Pairs CH3COOH(aq) +H2O(l) CH3COO–(aq)+ H3O+(aq) H+ conjugate base acid H+
Conjugate Acid/Base Pairs NH3 (aq) + H2O(l) NH4+(aq)+ OH–(aq) H+ H+ base conjugate acid